Langmuir 1993,9, 2330-2336
2330
Alcohol Aggregation at Hydrophobic Monolayer Surfaces and Its Effect on Interfacial Redox Chemistry Stephen E. Creager’ and Gary K. Rowe Department of Chemistry, Indiana University, Bloomington, Indiana 47405 Received May 3, 1993. I n Final Form: July 6,1993@ The aliphatic alcohols 1-butanol, 1-hexanol, 1-octanol, 1-decanol, and 2,2,4-trimethylpentanol were found to aggregate from aqueous solution onto the surface of hydrophobic monolayers of alkanethiols self-assembled onto gold electrodes. The aggregates were characterized via their effect on the redox properties of several ferrocene derivatives immobilized as minority components in the monolayers and on the interfacialcapacitance of the electrodes. Alcohol aggregationcausesa decrease in the overallcapacitance at the electrode and dramatically shifts the redox potential for ferrocene oxidation in a positive direction relative to the potential observed in the absence of alcohol. The magnitude of the shift in ferrocene redox potential is a function of the alcohol concentration in solution,reaching a maximum for solutions saturated with alcohol, and of chain length for the n-alcohols, reaching a maximum for 1-decanol(the longest alcohol tested). Aggregates form only when the monolayer presents a hydrophobic surface to the contacting solution;use of an alkanethiol with a hydroxyl terminal substituent or of a ferrocene derivative containing a quatehary ammonium group completely quenches the effect. A structural model is proposed in which aggregate layers formed from saturated alcohol solutionsare approximately1.1-1.5 nmthick and relatively disordered, with properties that are not greatly different from those of the bulk alcohol liquids.
Introduction Outer-sphere electron transfer is an example of a chemical reaction that is strongly affected by the properties of the microenvironment surrounding the reactants.l For example, the rate of outer-sphere electron transfer is determined in large part by the energy required to polarize the medium around the donor and acceptor to produce a transition-state geometry from which resonant electron transfer takes place? This energy is lowest when the medium is of low polarity, i.e. when the optical and static dielectric constants of the medium are of similar magnitude, as expressed mathematically in the celebrated Marcus theory of electron-transfer kinetic^.^ A potentially attractive research strategy for studying electron-transfer reactions is in fact to vary the nature of the medium in which the reaction takes place and then to correlate the rates with the properties of the medium.P12 Factors such as changes in reaction driving force, changes in electronic coupling between reactants, and solvent dynamical effects often have to be accounted for in such studies; however the study of solvent effects in electron transfer remains an active and prosperous endeavor. Microenvironmental effects in electron transfer are not limited to bimolecular reactions in homogeneous solution, however, and there have been several reports of solvent effects on outer-sphere electron transfer across the metal/ electrolyte solution interface.6Il2 Correlating the interfacial structure of electrode surfaces with observed patterns of
* Abstract published in Advance ACS Abstracts, August 15,1993.
(1) Cannon, R. D. Electron TransferReactiom; Butterworth London, 1980; pp 351. (2) Brunschwig, B. S.; Ehrenson, S.; Sutin, N. J. Phys.Chem. 1986,90, 3657. (3) Marcus, R. A.; Sutin, N. Biochim.Biophys. Acta 1985,811, 265. (4) Efrima, S.; Bixon, M. J. Chem. Phys. 1976, 64, 3639. (5) Fawcett, W. R.; Blum, L. Chem. Phys.Lett. 1991,187, 173. (6) Gennett, T.; M h e r , D. F.; Weaver, M. J. J. Phys.Chem. 1985,89, 2787. (7) Gochev, A.; McManis, G. E.; Weaver, M. J. J. Chem. Phys. 1989, 91,906. (8) Hupp, J. T.; Weaver, M. J. J. Phys.Chem. 1984,88, 1860. (9) Rips, I.; Jortner, J. J. Chem. Phys. 1987,87,6513. (10) Rips, I.; Jortner, J. J. Chem. Phys. 1987, 87, 2090. (11) Rips, I.;Jortner, J. Chem. Phys.Lett. 1987, 133, 411. (12) Weaver, M. J. J. Phys. Chem. 1979,83, 1748.
electrochemical reactivity is often difficult, however, even for relatively simple systems. Recently, progress has been made in this area by studying electron-transfer reactions of redox-active moleculesthat are held a t specific distances and orientations relative to an underlying electrode by virtue of being covalently bound to the outer portion of a semirigid self-assembled monolayer on the electrode surface. The most heavily studied system of this type is that of alkanethiols self-assembled onto gold? Examples of redox systems that have been immobilized onto gold using thischemistry include ferrocenes,1P29v i o l o g e n ~ , ~ * ~ ~ quin0nes,2~*~~ imide^:^^^^ and various ruthenium and osmium amine and polypyridyl complexes.3630 Several detailed studies have appeared that address the potential (13) Dubois, L. H.; Nuzzo, R. G. Annu. Rev. Phys.Chem. 1992,43,437. (14) Chidsey, C. E. D.; Bertozzi, C. R.; Putvinski, T. M.; Mujsce, A. M. J. Am. Chem. SOC. 1990,112,4301. (15) Chidsey, C. E. D. Science 1991,251,919. (16) Collard, D. M.; Fox, M. A. Langmuir 1991, 7,1192. (17) Curtin, L. S.; Peck, S. R.; Tender, L. M.; Murray, R. W.; Rowe, G. K.; Creager, S. E. Anal. Chem. 1993,65,386. (18) Rowe, G. K.; Creager, S. E. Langmuir 1991, 7, 2307. (19) Creager, S. E.; Rowe, G. K. Submitted for publication in J.
Electroanal. Chem. (20) DeLong, H. C.; Donehue, J. J.; Buttry, D. A. Langmuir 1991,7,
2196. (21) Duevel, R. V.; Corn, R. M. Anal. Chem. 1992,64,337. (22) Hick”, J. J.; Zou,C.; Ofer, D.; Harvey, P. D.; Wrighton, M. S. J. Am. Chem. SOC.1989,111, 7271. (23) Hickman,J.J.;Ofer,D.;Laibinis,P.;Whitesides,G.M.; Wrighton, M. S. Science 1991,252,688. (24) Hickman, J. J.; Ofer, D.; Zou, C.; Wrighton, M. S.; Laibinis, P. E.; Whitesides, G.M. J. Am. Chem. SOC.1991,113, 1128. (25) Popenoe, D. D.; Deinhammer, R. S.;Porter, M. D. Langmuir 1992, 8, 2521. (26) Tsuteumi, H.; Furumoto, S.; Morita, M.; Matauda, Y. J. Electrochem. SOC.1992,139, 1522. (27) Shimazu, K.; Yagi,I.; Sato, Y.; Uosaki, K. Langmuir 1992,8,1385. (28) Uosaki, K.; Sato, Y.;Kita, H. Electrochim. Acta 1991,36, 1799. (29) Uosaki, K.; Sato, Y.; Kita, H. Langmuir 1991, 7, 1510. (30) Delong, H. C.; Buttry, D. A. Langmuir 1992,8, 2491. (31) Lee, K. A. B. Langmuir 1990,6, 709. (32) Katz, E. Y.; Borovkov, V. V.; Evstigneeva, R. P. J. Electroanal. Chem. 1992,326, 197. (33) Kwan, W. S. V.; Penneau, J. F.; Miller, L. L. J.Electroanal. Chem. 1990,291, 295. (34) Kwan, W. S. V.; Atanasoska, L.; Miller, L. L. Langmoir 1991, 7 , 1419. (35) Finklea, H. 0.; Ravenscroft, M. S.; Snider, D. A. Langmuir 1993, 9, 223. (36) Finklea,H. 0.; Hanshew,D. D. J . Am. Chem. SOC. 1992,114,3173.
0743-7463/9312409-2330$04.00/0 0 1993 American Chemical Society
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Alcohol Aggregation at Hydrophobic Surfaces dependence of the electron-transfer rates in such syst e m ~ . Much ~ ~ of* our ~ own ~ ~work ~ in this area has focused on the importance of the local microenvironment at the electrode surface in determining the redox potential for immobilized redox-active groups in such layers; for monolayers in which ferrocenylalkanethiob were coadsorbed with long-chain alkanethiols, a local microenvironment of reduced polarity was invoked to explain the dramatic positive shifts that were observed in the ferrocene redox potential.18 Similar behavior has been reported for several related s y s t e m ~ . ~ ~ 3 ’ * ~ ~ In this paper we will show that the local microenvironment at an electrode surface can be strongly affected by the presence of aliphatic alcohols in the electrolyte solution that contacts the electrode. For hydrophobic monolayers of ferrocenylalkanethiols coadsorbed with alkanethiols, we show that the alcohols form surface aggregates that further reduce the polarity of the microenvironment around ferrocene, yielding a further positive shift in the ferrocene redox potential. The effect is operative only for monolayers with hydrophobic surfaces; monolayers prepared from functionalized alkanethiols that present hydrophilic surfacesto the solution are unaffected by alcohols in solution. A structural model for the alcohol aggregate layer on the monolayer surface is suggested.
Experimental Section Electrochemistry. Gold electrodeswere prepared by sealiig
gold wires in an insulating epoxy (EPON 825, Shell, cured with 1,3-phenylenediamine),followed by sanding and polishing to expose gold disks of macroscopic area 0.011 cm2(estimated from scanningelectronmicrographs). Polished electrodeswere briefly treated with 1:l aqua regia-water prior to coating.@ Selfassembled monolayers were prepared immediately after aqua regia treatment by exposing the treated electrode to a solution containing the thiol or thiols of interest for not less than 12 h. All coating solutions were made from ethanol except those containing (6-mercaptohexyl)(ferrocenylmethyl)dimethylammonium hexafluorophosphateand hexanethiol, which were made from acetonitrile. (Electrodescoated from this particular solution were subsequently soaked in an ethanol solution containing only hexanethiol; electrodes not subjected to this posttreatment typicallyshowed high background currentswhich tended to obscure the ferrocene electrochemistry.) Electrodes were rinsed with ethanol and water prior to use in electrochemical experiments. Electrochemical experiments were performed with a singlecompartment three-electrode cell that was continuously purged with water-saturated nitrogen. Electrochemical data were acquired using an EG&G PAR Model 362 scanning potentiostat. Capacitance data were obtained by superimposing a 20-Hz, 20mV (peak-to-peak)triangle wave onto the dc potential applied by the potentiostat to the working electrode;the resulting square wave current signal was monitored with a Tektronix Model 2430 oscilloscope, and the magnitude of the square wave was used to calculate the interfacial capacitance. Contact angles were measured using a h e - H a r t Model 100 contact angle goniometer. Gold electrodes for use in contact-angleexperimentswere thin-film electrodesprepared by vacuum evaporation onto glass. A (mercaptopropy1)triethoxysilanelayer was applied to the glass prior to gold coating to promote adhesion.“ Materials. Electrolytesolutionswere prepared from distilled/ deionized water (Barnstead Nanopure) using methanesulfonic acid (Aldrich). Electrolyte solutions containing alcohols were prepared by mixing the appropriate quantities of alcohol-free and alcohol-saturated electrolyte; this was done to avoid the measurement problem associated with preparing very dilute aqueous solutions of alcohols (for example, the solubility of (37) Redepenning, J.; Tunison, H. M.; Finklea, H. 0.Langmuir 1993, 9,1404. (38) Obeng, Y. S.; Bard, A. J. Langmuir 1991, 7, 195. (39) Acevedo, D.; Abruna, H. D. J. Phys. Chem. 1991,95,9590. (40) Creager, S. E.;Hockett, L. A.;Rowe, G . K. Langmuir 1992,8,854. (41) Goss, C. A.; Charych,D. H.; Majda, M. Anal. Chem. 1991,63,85.
8
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031 0
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% Saturation Figure 1. Effect of 1-octanolon the redox potential for oxidation
of ferrocenylhexanethiolimmobilized with n-hexanethiolon gold (rFc = 5 X mol cm-3. Redox potentials were estimated from cathodic cyclic voltammetric peak potentials; voltammograms were acquired in aqueous 0.10 M methanesulfonic acid. Splitting between anodic and cathodic peak potentials is always close to zero. Insert: voltammograms for octanol-freeelectrolyte (top) and octanol-saturated electrolyte (bottom). S = 10 nA, sweep rate = 100 mV 8-1.
1-octanol in water is 0.59 mg mL-l at 26 OC“) and to minimize the possibility that slow dissolution of alcohol in water would affect the meaeurements. Aliphaticalcohols1-butanol,1-hexanol, 1-decanol,and 2,2,4-trimethylpentanol(all from Aldrich), l-octan01 (Fisher), and alkanethiols n-hexanethiol, n-nonanethiol, and n-decanethiol(all from Aldrich)were reagent gradeand were usedwithout further purification exceptfor n-hexanethiol,which was passed through activated alumina prior to use. Tetradecanethiolwas prepared from n-bromotetradecaneby a literature methodq and was purified by Kugelrohr distillation prior to use. SMercaptohexanolwas prepared and purifiedin a similar manner using Sbromohexanol (Aldrich) as starting material. The preparation of the three ferrocene-containing alkanethiols is described elsewhere.l8
Results and Discussion Cyclic Voltammetry of Immobilized Ferrocene Groups. The major experimental finding of this work is that the presence of aliphatic alcohols, particularly n-alcohols, in an electrolyte solution can dramatically affect the redox properties of ferrocene immobilization in selfassembled monolayers. Specifically, the redox potential for ferrocene is observed to shift in a positive direction when such alcohols are present. Figure 1 illustrates this effect for a monolayer of ferrocenylhexanethiol coadsorbed with n-hexanethiol and then placed in contact with an aqueous electrolyte solution containing varying amounts of 1-octanol. The cyclic voltammograms in the insert correspond to electrolyte solutions containing no octanol (top)and a saturation level of octanol (bottom). It is clear that the presence of octanol has caused the redox potential to shift in a positive direction, in this case by a maximum of 320 mV. The shift is of thermodynamic and not kinetic origin, as indicated by the equivalent shifts in both the anodic and cathodic peak potential. The magnitude of the shift depends critically on the concentration of octanol (42) Stephen, H.; Stephen, T. Solubilities of Inorganic and Organic Compounds; MacMillan Co.: New York, 1963. (43) Cossir,B. C.; Fournier,J. 0.; Fields, D. L.; Reynolds, D. D. J. Org. Chem. 1962,27,93.
2332 Langmuir, Vol. 9, No.9, 1993
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Creager and Rowe
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E (V vs. Ag/AgCI)
Figure 2. Effect of 1-octanol on the oxidation of ferrocenylhexanethiol immobilized with 6-mercaptohexanol on gold (rFc = 1.0 x mol cm-l): (A) electrolyte solution (aqueous 0.1 M methanesulfonic acid) without octanol;(B)electrolyte solution saturated with 1-octanol. 5' = 10 nA,sweep rate = 100 mV s-l. 0.0
in solution. This is illustrated graphically in Figure 1, which shows that the maximum shift in redox potential is approached monotonically with increasing octanol concentration, the relative shift being greatest at lower octanol concentrations and approaching a constant value at saturation. A detailed description of the proposed cause of this shift in redox potential is given below; however in general terms we propose that the octanol aggregates onto the surface of the self-assembled monolayer, thereby creating a less polar local microenvironment that favors ferrocene relative to ferricinium and shiftsthe equilibrium position of the surface redox reaction. This explanation is qualitatively similar to one that we invoked in our earlier work on shifts in redox potential for ferrocenylhexanethiol coadsorbed with n-alkanethiols of variable chain length,4 the difference being that in the earlier work there were no alcohols present in solution and the changes in local microenvironment were due to structural constraints within the monolayer itself. A second possibility, also discussed below, is that the alcohol layer may generate an unusual ion and/or potential distribution that causes an apparent redox potential shift. One advantage of the self-assembly method aa a means of studying electron transfer is that one can introduce systematic changes into the structure of the monolayersolution interface by using functionalized alkanethiols to prepare various mixed monolayers. We have adopted this strategy here as a way of probing how the shift in redox potential induced by alcohols in solution depends on the details of the interfacial structure at the monolayersolution interface. Specifically, we have varied the terminal functionality on the coadsorbate alkanethiol, with the goal of introducing polar substituents that should render the surface less hydrophobic and therefore perhaps less likely to interact with alcohols in solution. Figure 2 illustrates how the voltammetry for a monolayer of ferrocenylhexanethiol coadeorbed with an n-alkanethiol with a hydroxyl group in the terminal position is affected by the presence of alcoholsin solution. Inspection of Figure 2 reveals immediately that there is now absolutely no effect of 1-octanol on the voltammetry of the immobilized ferrocene, even when the octanol is present at the saturation level. The effect that caused the dramatic shifts evident in Figure 1 is completely quenched when the monolayer is rendered hydrophilicby using an alkanethiol coadsorbate with a polar terminal substituent. Another way of introducing a structural perturbation into the monolayer-solution interface is to use ferrocenylalkanethiols in which the functional group linking (44)Creager, S. E.;Rowe, G.K.Anal. Chim. Acta 1991,246, 233.
0.2
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% Saturation Figure 3. Effect of 1-octanolon the redox potential for oxidation ofN-(7-mercaptoheptyl)ferrocenecarboddeimmobilized with n-nonanethiol on gold (rF, = 6 X It1*mol cm-l): triangles; data for N-(7-mercaptoheptyl)ferrocenecarboxamide;circles, data from Figure 1. Other conditions are identical to those in Figure 1. Insert: voltammograms for octanol-freeelectrolyte (top) and octanol-saturated electrolyte (bottom). T = 10 nA.
ferrocene to the alkanethiol chain is varied. We recently prepared such a set of compounds and have studied ita behavior as self-assembledmonolayers with n-alkanethiols on gold.1° In the context of the present work we are interestad primarily in whether the nature of the functional group linking ferrocene to the alkanethiol chain has any effect on the ability of alcohols in solution to perturb the ferrocene redox potential. Figure 3 illustrates what happens when the relatively nonpolar ferrocenylhexanethiol is replaced by a more polar molecule, specifically N-(7-mercaptoheptyl)ferrocenecarboxamide, in the mono. layer. The coadsorbate alkanethiol was in this case n-nonanethiol, which was chosen because the alkyl chain in nonanethiol has the same number of atoms aa does the chain linking ferrocene to the electrode surface (if the two "links" in the carbosamide linking group are counted). It is clear that substitution of the carbosamide group for the methylene group has had little effect on the ability of 1-octanolto induce a shift in redox potential; in fact, the curves showing redox potential vs octanol concentration for the two monolayers in question are nearly parallel, the offset presumably being caused simply by the different intrinsic redox potentials for the two ferrocenederivatives. Figure 4 illustrates what happens when an even more polar substituent, specifically a cationic quaternary ammonium group, is used to link ferrocene to the alkanethiol chain. Adding 1-octanol to the electrolyte solution now has only a very minor effect on the redox potential for ferrocene oxidation, even though the coadsorbate is again n-hexanethiol which might be anticipated to present the same hydrophobic surface to the solution as waa present in previous experiments. Apparently the interfacial structural perturbation introduced by the positivelycharged quaternary ammonium group is so great aa to disrupt any local structures that might form as a result of alcohols interacting with the monolayer surface. The local structure at the monolayer surface is dominated by the requirement to solvate the ion (or ion pair) that resides there. We have explored several other perturbations of monolayer structure, mostly involving variations in relative chain
Langmuir, Vol. 9, No.9, 1993 2333
Alcohol Aggregation at Hydrophobic Surfaces
Table I. Effect of n-Alcohols on the Redox Potential of Immobilized Ferrocenes monolayer structureb (Fc)C&H /C&He (Fc)C&H/CsSH (Fc)C&H/CioSH (Fc)C&H/HOC&H (Fc)C&H/C&H
alcoholc CsOH CsOH CsOH CsOH CiOH CeOH CsOH i-CsOHf CioOH CsOH CsOH CsOH CsOH
(Fc)CONHC~SH/C&H (Fc)CONHC~SH/C&H (Fc)CONHC,SH/Ci&H (FC)CH~N+(CH&C&H/C~SH
Eo (no alc) (Vd) +0.33 +0.41 +0.66 +0.27 +0.43 +0.43 +0.43 +0.42 +0.43 +0.52
E O
(W/al@)(V) +0.60 +0.73 +0.81 +0.27 +0.58 +0.59 +0.69 +0.66 +0.72 +0.67
+0.68
+0.84
+0.81 +0.63
+0.95 +0.69
AE (+ale) (V) +0.27 +0.32 +0.15 0.00 +0.15 +0.16 +0.26 +0.24 +0.29 +0.15 +0.26 +0.14
+OM
Alcohol present at saturation level. Nomenclature: (Fc)CaSH 3 6-ferrocenylhexanethiok(Fc)CONHC7SH = N(7-mercaptohepty1)ferrocenecarboxamide; (Fc)CH2N+(CH&C&H = (ferrocenylmethyl)(6-mercaptohexy1)dimethyla"onium ion; CnSH = n-alkanethiol of chain length n. Nomenclature: CnOH = n-alcohol of chain length n. Dielectric constanta of bulk alcohols are ae follows: 1-butanol, 17.8; 1-hexanol, 13.3; 1-octanol, 10.3; 1-decanol, 8.1. Volta vs Ag/AgCVsat'd KC1. e Coverage of ferrocene is always less that 1 X 10-11 mol cm-2. f i-CeOH = 2,2,4trimethylpentanol.
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Figure 4. Effect of 1-octanol on the oxidation of (ferrocenylmethyl)(6-mercaptohexyl)dimethyla"onium ion (initially present as the hexafluorophosphate salt) immobilized with mol cm-l): (A) electrolyte hexanethiol on gold (rpc = 7 X solution (aqueous0.1M methanesulfonicacid) without 1-octanol; (B) electrolyte solution saturated with 1-octanol. R = 25 nA, sweep rate 100 mV s-l.
length of the two alkanethiols that comprise the monolayers. The results of these experiments are summarized in Table I in the form of redox potentials for ferrocene oxidation in each monolayer in two different electrolyte solutions, one without alcohol and one saturated with alcohol. In every case except that of ferrocenylhexanethiol coadsorbed with 6-mercaptohexanol, the presence of alcohol in solution causes the redox potential for ferrocene oxidation to shift in a positive direction. The column on the far right of Table I shows the difference in ferrocene redox potential for the alcohol-free and alcohol-saturated solutions. Within the ferrocenylhexanethiol/alkanethiol series, the effect of 1-octanol, a representative alcohol, appears to be greatest when the length of the coadsorbate alkanethiol chain is equal to or lees than that of the alkyl chain that is attached to ferrocene. A similar effect is observed in the ferrocenecarboxamideseries, except that the shift is greatest when the length of the coadsorbate alkanethiol chain exactly matches that of the chain linked to ferrocene. Of course, monolayers formed using coadsorbate alkanethiols that are significantlylonger than the chains attached to ferrocene always show large positive shifta in ferrocene redox potential relative to monolayers formed using shorter coadsorbates, as discussed in our earlier publication.44 The point here is that while the overall effect on the redox potential becomes greater as the chain length of the coadsorbate alkanethiol increases, the relative effect of alcohols in solution is greatest when the coadsorbate alkanethiol chain is of equal or shorter length relative to that of the chain attached to ferrocene.
This presumably occurs because when the coadsorbate alkanethiol chain is long, the ferrocene is already in a relatively nonpolar environment so that further changes in microenvironmentbrought about by the alcohols layer are of lesser importance. The effect described above is not limited to 1-octanol and in fact is observed in varying degrees for each of the series of n-alcoholsfrom butanol to decanol (alcoholsbelow butanol are miscible with water, and those above decanol are so nearly insoluble as to make experiments difficult), as well as for the branched octanol isomer 2,2,4-trimethylpentanol. Table I includes data on the shift in redox potential for a particular ferrocenylhexanethiol/hexanethiol monolayer in contact with saturated aqueous solutions of each of these alcohols. Inspection of the data reveals that the redox potential in saturated alcohol solutions shifts to progressively more positive potentials as the chain length of the alcohol increases, reaching a maximum for a solution saturated with 1-decanol. (Note that the redox potential for the indicated layer in an octanol-saturatedsolution on this specific electrode differs slightly from that reported earlier in the table for a nominally identical layer under similar conditions. This difference reflects the level of experiment-to-experiment uncertainty in these measurements. Both values are included in the table to give an approximate indication of the overall uncertainty in the data; uncertainties in peak potential for a given electrode are always less and are in the range of 10 mV.1 The shift in redox potential for the different n-alcohols follows the progressive decrease in dielectric constants for the n-alcohols as bulk liquids (footnote c in Table I), which suggests that the changes in ferrocene microenvironment caused by n-alcohol aggregation reflect at least in part the bulk properties of the n-alcohols. The physical model we propose to explain these observations is that of an aggregate layer of alcohols on the monolayer surface, the formation of which is driven by ti favorable hydrophobic interaction between the monolayer surface and the alkyl chains on the alcohols. Two further pieces of data which support this model are as follow: (i) the presence of n-alcohols in solution has absolutely no effect on the electrochemistry of an immobilized ferrocenylalkanethiol when it is studied in a nonaqueous electrolyte solution such as acetonitrile, and (ii) the presence of n-alcohols in solution has no effect on the redox potential of water-soluble ferrocenes such as hydroxymethylferrocene. Long-chain alcohols are known to be surface-active in so it seems plausible that (45) Hommelen, J.
R.J. Colloid Sci.
19S9, 14, 386.
2334 Langmuir, Vol. 9, No. 9,1993
Creager and Rowe
Table 11. Effect of n-Alcohols on the Capacitance of Monolayer-Coated Electrodes. monolayer structure (Fc)CeSH/C4SH (Fc)C&H/C&H (FdCeSH/CioSH (Fc)C&H/HOC&H (Fc)CeSH/CeSH
(Fc)CONHC,SH/CeSH (Fc)CONHC,SH/CgSH (Fc)CONHC,SH/C~~SH (Fc)CHzN+(CHs)&eSH/C&H (no ferrocene)/C&H
alcohol CsOH CeOH CsOH CsOH CiOH CeOH CsOH i-CeOH CioOH CsOH CsOH CsOH CsOH CsOH
CDL(no alc) (pF cm-2 5.6 3.7 2.4 1.5 3.8 3.8 3.8 3.8 3.8 3.6 2.4 2.0 1.7
3.8
CDL( w / ~ c ) (pF cm-2) 2.9 2.5 1.9 7.5 3.0 2.8 2.5 2.1
2.1 2.5 1.9 1.6 6.1 2.5
CWWltOb (pF cm-2) 6.0 1.7 9.1 14.3 10.6 1.3 9.3 4.1 8.2 9.1 8.0 (52) 1.3
0 Nomenclature ie identical t o that in Table I. Capacitance calculated from CDLmeasured with and without alcohol, as described in the text. C Electrode area taken as 0.011 cm2.
they would tend to aggregate at a hydrophobic solid surface in water. Similar aggregation phenomena have in fact been reported by Miller and co-workers in their work on the effect of alcohols on the lateral mobility of redoxactive surfactants on metal oxide surfaces coated with organosilanesa and by Barton and co-workersin their work on alcohol aggregation in surfactant adlayers on liquid chromatographystationary phases.47 It furthermore seems reasonable that such an aggregate layer could change the local microenvironment a t an electrode surface and that the changed microenvironment would manifest itself as a shift in redox potential for an immobilized redox-active molecule. Physical models have been described that predict such shifts in redox potential as the solvent surrounding a redox-active molecule is changed.& Close inspection of the voltammograms in the insert of Figure 1reveals that the shape of the voltammetric wave is more ideal, i.e. the peaks are more narrow, in the presence of 1-octanol than in its absence. This effect is seen even at subsaturation levels of 1-octanol;for example, the fullwidth at half-maximum of the ferrocene wave in a ferrocenylhexanethioVhexanethio1monolayer is typically between 150 and 160 mV in the absence of 1-octanol but is reduced to 95-105 mV with only 20% of the saturation level of 1-octanol present in solution. We speculate that this reflects a greater homogeneity in the range of microenvironments experienced by the ensemble of ferrocenes in the monolayer. This speculation implicitly suggests that the cause of the unusually broad waves observed in the absence of alcohol is that there is initially a range of redox potentials for different ferrocene sites in the monolayer. Alternative causes of broadening such as repulsive interactions among oxidized and/or reduced site* are thereby implicitly considered to be less important. InterfacialCapacitanceMeasurements. In addition to the voltammetric data presented and discussed above, we have also made measurements of the interfacial capacitanceof monolayer-coated electrodesin the presence and absence of alcohols. Capacitance was measured at a potential far from the ferrocene redox potential, typically at 0.00 V vs AgiAgCl, to avoid complications from current associated with faradaic electrochemistry. The results of these measurements are summarized in Table I1for several different monolayers contacting solutionswith and without alcohols present. The general trend is that having an (46) Miller, C. J.; Widrig, C. A.; Charych, D. H.; Majda, M. J. Phys. Chem. 1988,92, 1928. (47) Barton,J. W.; Fitzgerald, T. P.; Lee, C.; ORear, E. A.; Harwell, J. H. Sep. Sci. Technol. 1988,23,637. (48)Strehlow, H. In The Chemistry of Non-Aqueous Solvents; Lagowski, J. J., Ed.;Academic Press: New York, 1966; Vol. I, p 129. (49) Brown, A. P.; Anson, F. C. A m l . Chem. 1977,49,1589.
alcohol present in solution causes the interfacial capacitance to be significantly, but not dramatically, smaller than it is in the absence of alcohol. The total interfacial capacitance for these coated electrodes can be thought of as being comprised of several individual capacitances, two of which correspond to the monolayer on the electrode surface (one for the alkanethiol layer and one for the aggregate alcohol layer) and one of which corresponds to the diffuse layer in solution. If one assumes that the measured capacitance in the absence of alcohol is equal to the capacitance of the alkanethiol monolayer (i.e., assume that the diffuse layer capacitance is sufficiently large that it can be ignored), and that the effect of the alcohol is simply that of an additional capacitance taken in series, then it is possible to estimate the capacitance of the alcohol aggregate layer by noting that the reciprocal of the measured capacitance in the presence of alcohol is the s u m of the reciprocals of the measured capacitance in the absence of alcoholand of the capacitance of the alcohol aggregate layer. Values for the capacitance of the alcohol aggregatelayer obtained in this way for several different monolayers and alcohols are included in the column on the right of Table 11. While there is some scatter in the data, there are also some clear trends. One such trend is that the capacitance of the alcohol aggregate layer varies fairly strongly with n-alcohol chain length, becoming smaller as chain length increases (entries 5,6,7, and 9 in Table 11). Another trend is that the aggregate layer capacitance for a given alcohol varies only weakly with the specific structure of the underlying monolayer, for example, the average capacitance for a 1-octanol layer is 7.8 f 1.0 pF cm-2 atop monolayers of alkanethiols from 4 to 14 carbons in length (entries 1, 2, 3, 7, 10, 11, 12, and 14 in Table 11). These generalizations are of course true only for monolayers that are essentiallyhydrophobic,since alcoholsdo not aggregate onto monolayers that are hydrophilic. The trend with respect to alcohol chain length is again consistent with the known dielectric properties of the bulk n-alcohol liquids; longer-chain n-alcohols have lower dielectric constants and therefore might be anticipated to produce aggregate layers of lower capacitance. The capacitance of the alcohol aggregate layers also follows in general terms the induced shifts in ferrocene redox potential; the layer with the lowest capacitance (decanol) also produced the largest shift in redox potential. This further reinforces the postulate that the properties of the alcohol aggregate layer reflect the properties of the liquid n-alcohols from which they are formed. Structure of the Aggregate Alcohol Layer. Having suggested that alcohols form aggregate layers on the
Langmuir, VoZ.9, No.9,1993 2335
Alcohol Aggregation at Hydrophobic Surfaces chart I
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A
B
C
surface of hydrophobic monolayers, we next seek to understand in more detail the structure of these aggregates. Nearly every possible structure that one could imagine for such an aggregatelayer atop a self-assembledmonolayer is strictly a bilayer; however, one can envision several limiting structures for such a bilayer including: (A) an organized second layer with the alkyl chains of alcohol molecules oriented approximatelynormal to the surface; (B) a loose, disorganized second layer with no preferential orientation of alcohol molecules; and (C) an organized but relatively thin second layer with the alkyl chains of alcohol molecules oriented approximatelyparallel to the surface. These limiting structures are illustrated in Chart I for an octanol layer. The underlying self-assembled layer is shown as being completely ordered; this approximation should perhaps not be taken too literally, though the alltrans, solid-likeconfiguration is well supported by Raman spectroscopic data for alkanethiols greater than five carbons in length on gold.50 Aggregate layers of variable thickness and with different structures at different depths are also possible, as are structures with some degree of interpenetration of the alcohol and alkanethiol layers. Distinguishing among these interfacial structures can be definitively accomplished only by using data from structure- and orientation-sensitive measurement tools, for example reflectance infrared spectroscopy or Raman spectroscopy. Such measurementswill be reported in due course; however in their absence it is still possible to make inferences about the interfacial structure by undertaking a more detailed analysis of the electrochemical data. C = eedd
(1)
One such analysis involves the magnitude of the capacitance of the alcohol aggregate layers. A simple expression that allows one to estimate the capacitance of a thin layer of a material of known thickness and dielectric properties is the Helmholtz equation (eq l),which states that the capacitance per unit area is given by the permittivity of the layer (the product of the dielectric constant of the layer and the permittivity of free space) divided by its thickness. Hence if one can estimate the dielectric constant for the aggregate layer, it would be possible to estimate the layer's thickness from the measured capacitance. Using a value of 7.8 X le2 F m-2 for the mean capacitance of l-octanol aggregate layer and a value of 10.3 for the dielectric constant of bulk l-octanol, we readily estimate the thickness of the l-octanol aggregate layer to be approximately 1.2 nm. Similar calculations for each of the series of n-alcohols consistently yield estimates of the aggregate layer thickness at saturation of between 1.1and 1.5 nm. These estimates depend critically on the value used for the dielectricconstant of the aggregate layer; our use of the values for the bulk alcohols assumes that the physical properties of the aggregate layer are not greatly different from those of the bulk alcohols. If we instead use a value of 2.3 for the dielectric constant (ie.,
the value for polyethylenes1), we can readily estimate the thickness of the l-octanol aggregate layer to be 0.26 nm. This thickness estimate is consistent with structure C but not with structure A in Chart I. Of course, the use of such a low dielectric constant assumes that the hydroxyl group in the alcohol layer has essentially no effect on the layer's dielectricproperties; if this estimate of the layer thickness is to be supported, then it becomes necessary to justify this assumption. It is instructive to consider whether the capacitance changes could be caused simply by intercalation of alcohols into the alkanethiol monolayer. Such intercalation phenomena have been observed for alcohols in surfactant layer^.^,^^ Applying eq 1to the capacitance measured for a pure hexanethiol monolayer in the absence of alcohol (entry 14in Table 11),and assuming a monolayer thickness of 0.66 nm (0.11 nm per CH2 group calculated assuming an average tilt angle of 30" off norma151),we estimate the effective dielectric constant of the monolayer to be 2.8, in excellent agreement with the estimate of 2.6 cited by Porter and co-workers from similar data for a series of alkanethiols.51 Minor differences could be due to a small population of defects, or to inadequacies in our estimate of the microscopic electrode area. Similar treatment of the data obtained in saturated l-octanol solution yields an effective dielectric constant of 1.9 for the surface layer if the thickness is assumed to remain constant. This value is clearly too low to be physically realistic, indicating that the alcohol must increase the thickness of the overall surface layer at least somewhat. This is not unexpected given the fundamental structural differences between a self-assembled monolayer and a surfactant layer; surfactant molecules are free to rearrange themselves to make room for intercalating molecules, however the components of a self-assembled monolayer are not, since they are both tightly packed and strongly held onto the surface of the solid substrate. Another way of analyzing the capacitance data is to compare the capacitance of the alcohol aggregate layer to that of a self-assembled alkanethiol monolayer film of similar composition, Le. an alkanethiol with a terminal hydroxyl group. Such a self-assembledmonolayer should have a structure very much like that in structure A of Chart I. For example, our value for the capacitance of a self-assembled monolayer of 6-mercaptohexanol is reported in Table I1as 7.5 pF cm-2. This value is somewhat higher than that obtained for a hexanethiol monolayer, which may indicate that the hydroxyl group in the 6-mercaptohexanol layer can make a dipolar contribution to the capacitance. Alternatively, the 6-mercaptohexanol monolayer may simply allow water and/or ions to penetrate more deeply, thereby reducing the effective monolayer thickness. In any case, the capacitance of an aggregate layer of hexanol atop a hexanethiol monolayer is 10.6 pF cm-l, which is clearly higher than the value for a 6-mercaptohexanol monolayer. This indicates that the hexanol aggregate layer is either thinner or has a higher effective dielectric constant than does the 6-mercaptohexanol monolayer, both of which effectively rule out structure A as it is drawn in Chart I. Another critical assumption in ths treatment is that the alcohol aggregate layer contains little or no ions. This is almost certainly not true at potentials positive of Eo where the ferrocene groups in the monolayer are in the ferricinium form; however it may well be true at 0.00 V vs Ag/AgCl. Partition coefficients for some simple binary electrolytes (HC1, LiC1, NaC1, and KC1) have been reported for the ~~~
(50) Bryant, M. A.; Pemberton, J. E. J. Am. Chem. SOC.1991,113, 82M.
(51) Porter, M. D.; Bright, T. B.; Allara,D. L.; Chidsey, C. E. D. J. Am. Chem. SOC.1987,109,3559.
2336 Langmuir, Vol. 9, No. 9, 1993 water-octanol system; as expected, the salts remain primarily in the aqueous phase.62 Even when salts partition into the octanol phase, they are reported to be highly associated; for example, the conductivity of l-octan01 in contact with 0.1 M HCl is reported to be approximately 1 X 1V S cm-l, which is 4 orders of magnitude less than the typical value for a binary electrolyte such as HC1 in water.62 We are not aware of any similar data for methanesulfonic acid; however we suppose that the situation is similar. We have also measured the contact angles of water droplets on these surfaces as a way of learning about the hydrophobicity of the various monolayers in the presence and absence of alcohol. It was thought that if structure A in Chart I is correct, then the contact angle in the presence of alcohol should drop to nearly zero because the monolayer surface would have become very hydrophilic as a result of it having exposed hydroxyl groups. Indeed it has been reported by several groups that surfacescoated with alkanethiols with a polar terminal substituent such as a hydroxyl group are completely wet by water, which reflects the relatively high surface tension of such a surface.63fMMeasuring contact angles in the presence and absence of alcohol is unfortunately complicated by the fact that the surface tension of water is iteelf dramatically changed by the alcohol;46hence changes in contact angle are expected even if the alcohol has no effect on the monolayer surface structure. Still,important insights can be obtained by measuring contact angles for different surfaces with water in the presence of alcohol. For example, a surface coated with a monolayer of n-hexanethiol exhibits a contact angle of 5 4 O when in contact witha droplet of electrolytesolution that is saturated with 1-octanol. A surface coated with a monolayer of 6-mercaptohexanol exhibits a contact angle of (i.e., very close to zero) when in contact with a similar droplet, also saturated with 1-octanol. We conclude from these data that the surface of the hexanethiol monolayer, even when it is covered by an octanol aggregate layer, does not resemble a close-packed layer of hydroxyl groups such as is present at the surface of the 6-mercaptohexanol layer. This again argues against a structure such as that in structure A in Chart I. The combined weight of all these data leads us to posulate that a structure similar to B in Chart I is most appropriate for describingthese alcohol aggregateslayers. Structure A is inappropriate given the capacitance and contact angle data presented above. Also, such a structure would be quite unlikely for the branched alcohol 2,2,4trimethylpentanol, yet the behavior of this alcohol is not greatly different from that of 1-octanol. Structure C is inappropriate since the dielectric constant required of the aggregate layer to produce the observed changes in capacitance is unreasonably low. Also, it is difficult to imagine how such a thin layer could exert such a strong influence on the redox properties of an immobilized ferrocene, particularly when that ferrocene is held to the surface by a long chain relative to the coadsorbate alkanethiol chain. Structures consisting exclusively of (52) Weatall, J. C.; Johnson,C. A.; Zhang, W. Enuiron. Sci. Technol. 1990,24,1803. (53) Folkera, J. P.; Laibinis, P. E.; Whitesidea, G.M. Langmuir 1992, 8, 1330. (54) Ulman, A.; Evans, 5. D.; Shnidman, Y.; Sharma, R.; Eilera, J. E.; Chang, J. C. J. Am. Chem. SOC.1991,113, 1499.
Creager and Rowe intercalated alcohols are also inappropriate, as discussed earlier. Structure B is consistent with the observation that the polarity of the local environment surrounding ferrocene seems to scale with the dielectric properties of the bulk alcohols. Of course, structures intermediate between those shown in Chart I, for example, with partial ordering of alcohol molecules or with different structures in different regions of the layer, are also possible, more specific statements regarding the detailed structure must await the availability of more in situ structure-sensitive surface characterization data. Finally, it is appropriate to again comment on the importance of the counterion distribution at the electrode surface. Counterion motion is known to be critically important in the electrochemistry of redox polymer films,ss*s8 and counterion distribution has been cited as a critical factor in determining both the thermodynamics and kinetics of redox reactions in self-assembled monol a y e r ~ . ~We ' ~ ~have suggested that free ions are unlikely to be present in an aggregate layer if its properties are similar to those of the bulk alcohol. One consequence of this is that the applied potential will not be completely dropped in the space between the ferrocene moieties and the electrode and that the ferrocenes will reside in a region of relatively high electric field. Apparent redox potentials that are shifted relative to values in the absence of such effects are predicted for such a situation,S7in agreement with what we observe. Another effect that is predicted for this situation is that the voltammetric waves will be significantlybroadened, we did not observe this, but rather observed the opposite effect of narrower waves in the presence of alcohol. In fact, the present situation is more complex in that ions will almost certainly be able to penetrate the alcohollayer once ferriciniumions are formed since the aggregate layer is undoubtedly quite fluctional and the driving force for counterion penetration would be strong. It furthermore seems likely that ion penetration into the aggregate layer would be a strong function of potential even in the absence of faradaic electrochemistry; at potentials near the potential of zero charge (pzc) there is little drivingforce for ions to penetrate the layer, however at other potentials the strong electric field across the layer may cause it to collapse and allow ions to get closer to the underlying metal electrode. Similar behavior has been reported for mercury electrodes in contact with aqueous solutions containing molecules such as n - a l c o h ~ l s .De~~ convoluting the relative importance of the polarity of the local microenvironment and of these unusual interfacial ion and/or potential distributions will therefore not be easy and will almost certainly depend on the availability of better in situ surface structural characterization data obtained as a function of applied potential.
Acknowledgment. This work was supported in part by a grant from the National Science Foundation (CHE9216361) and by a summer fellowship to G.K.R.from the Analytical Division of the American Chemical Society. (65) Majda, M. In Molecular Design of Electrode Surjaces, Murray, R. W., Ed.;John Wdey & Sone, Inc: New York, 1992; Vol. 22, p 427. (56)Jernigan, J. C.; Murray, R. W. J. Phys. Chem. 1987,91,2030. (57) Smith, C. P.; White, H. S . Anal. Chem. 1992,64,2398. (58) Creager, S. E.; Weber, K. Langmuir 1993,9,&14. (59) Gileadi, E.; Kirowa-Eianer,E.; Penciner, J. In Interfacial Electrochemistry; An Experimental Approach; Addison-Wesley Publishing Co.: Reading, MA, 1975; p 41.
